According to our theory atomic size should increase in groups from top to bottom but in group 13 Aluminium has atomic
radius 143 pm and Gallium has atomic radius 135 pm which is less WHY??????
The reason for the anomaly in atomic radii between aluminum and gallium in Group 13 lies in the electron configuration and the shielding effect.
In a group, as you move down, a new energy level (shell) is added, resulting in an increased atomic size. This is because there are more energy levels to accommodate more electrons, leading to a larger atomic radius.
However, in the case of aluminum (Al) and gallium (Ga), the trend is reversed due to the presence of the d-block transition metals in between.
Aluminum (Al) has the electron configuration [Ne]3s^23p^1, and gallium (Ga) has the electron configuration [Ar]3d^104s^24p^1.
The d-block transition metals have partially filled d orbitals, which are located closer to the nucleus compared to the s and p orbitals. These d electrons experience a greater nuclear attraction, resulting in a contraction of the atomic size.
Therefore, despite being in the same group, gallium (Ga) has a smaller atomic radius compared to aluminum (Al) due to the presence of the d-block transition metals and the shielding effect caused by the d-electrons.
The atomic size or atomic radius is generally expected to increase from top to bottom within a group on the periodic table. However, there are certain exceptions or anomalies that can occur due to unique electron configurations and other factors.
In the case of group 13, which is also known as the boron group, the elements in this group have a general trend of increasing atomic size from Boron to Thallium. Boron has a smaller atomic radius compared to the elements below it in the group.
Now, when we specifically consider the difference in atomic radii between Aluminum (Al) and Gallium (Ga), we observe that Al has a larger atomic radius (143 pm) compared to Ga (135 pm). This contradicts the trend of increasing atomic size from top to bottom within the group.
To understand this discrepancy, we need to consider the electronic configurations of Aluminum and Gallium. Aluminum has an electron configuration of [Ne] 3s² 3p¹, while Gallium has an electron configuration of [Ar] 3d¹⁰ 4s² 4p¹.
The presence of the extra d-electrons in Gallium causes more effective nuclear charge or attraction towards the outermost electrons. This results in a greater pull on the outermost electron cloud, leading to the contraction of the atomic radius in Gallium compared to Aluminum.
In general, it is important to keep in mind that while there are predictable trends in atomic size on the periodic table, individual elements and their unique electron configurations can sometimes lead to exceptions like this one.